Fundamentals of Inorganic Chemistry

Inorganic Chemistry Topics discussed:

1. Atomic Structure:
2. Quantum numbers
3. Electronic configuration:
4. Periodic Table:
5. Determine the group and period of an element with atomic number
6. Periodic Table Trends (periodic table properties)

See Also:  Groups of Periodic Table and Trends in Periodic Table

Atomic Structue:

-         Atom consist of a nucleus (composed of protons and neutrons) surrounded by electrons.

-         Since atom is neutral the number of positive protons = the number of negative elctrons.

-         The elements symbols can be written as follow:

                                        _{Atomic number      z}X^{y      Mass number}

 

-         Definition of Atomic number: is the number of prtons in the nucleus OR the number of electrons around the nucleus.

-         Definition of Mass number: is the sum of the number of protons and neutrons.

-        Definition of Isotopes: Atoms that have the same atomic number but different in the mass number.

-         Examples of isotopes:

                                      ₆C¹²     ₆C¹³    ₆C¹⁴

 

-        Relative atomic mass: Samples of elements contain mixture of isotopes; their proportions are given by the abundance.

    Ex:   If the abundance of hydrogen isotopes were ₁H¹=70% - ₁H²=20%    ₁H³=10%

    Then the relative atomic mass of hydrogen =70x1/100 + 20x2/100 + 10x3/100 =1.4

 

-       Dual nature of electron: Electrons behave as both waves (electromagnetic waves) and particles.

Electromagnetic waves: are the waves that have both electrical and magnetic properties.

The energy of electromagnetic waves can be described using frequency (ν) and wavelength (λ):

                                              E= hν=hC/ λ       

                                               Also, E=mC²

                                           Then, hC/ λ= mC²

                                                   λ =h/mC

                                      E=Energy.                   m=mass.

                              h=Plank^,s  constant.        C= Velocity of light.

 

-        Uncertainty principle: this principle states that we cannot precisely measure both the location and the velocity of electron.

So, we talk about the probability of finding the electron around the nucleus,this is what is called “electron cloud”.

 

Quantum numbers: the numbers which determine the electronic configuration of  atoms.

 

(a)- Principle quantum number(n):

        - Describe the energy of states.

        - Take intger numbers : 1,2,3,4……

        - As (n)  increase the electron energy increase and the electron become less bounded to the nucleus .

         - The numbers take the following letters:

                     1=K  ,   2= L   ,  3=M    , 4=N  ………………….

 

 (b)- Angular quantum number(   ):

        - Determine the number of sub-levels.

        - Define the shape of the orbtals.

        - Take intger numbers from 0 to (n-1).

        - The numbers take the following letters

                 0=s      ,    1=p    ,    2=d     ,     3=f     ,    4=g

 

(c)- Magnetic quantum number(m  ):

        - Determine the number of orbitals in the sub- level.

        - Take integer numbers from        to         including  “0”.

 

(d)- Spin quantum number(m_s):

        - Determine the magnetic field result from the spin of electron

        - Take the values  + ½  or  - ½.

 

Examples for quantum numbers

 

Symbol of main level	n	l	m	Number of sub levels	Symbol of sub level	Number of orbitals
K	1	0	0	1	1s	1
L	2	0 1	0 +1,0,-1	2	2s, 2p	1 3
M	3	0 1 2	0 +1,0,-1 +2,+1,0,-1,-2	3	3s, 3p 3d	1 3 5
N	4	0 1 2 3	0 +1,0,-1 +2,+1,0,-1,-2 +3,+2,+1,0,-1,-2,-3	4	4s 4p 4d 4f	1 3 5 7

          

          

-         Orbitals:

-The electron capacity of each orbital is two electrons.

-The shapes of the  orbitals in each sub level are given as follow:

(a)- S sublevel:

       Consist of one orbital with spherical shape as illustrated in the figure:

(b)- P sublevel:

        Consist of three orbitals P_x , P_y , and P_z.  and their shapes are as follow:

 

 

(c )- d  sublevel:

        Consist of five orbitals d_xy , d_yz , d_xz  , dx²-y²  , and  dz² 

(d)- f sublevel:

             consist of seven orbitals with very complicated shapes.

 

Electronic configuration:

-         The energy levels are different in energy and the energy of the level increase by increasing the distance from the nucleus

-         The arrangement of the sulevels is as follow:

    1s  2s  2p  3s  3p  4s  3d  4p  5s  4d  5p  6s  4f  5d  6p  7s   5f  6d  7p

 

-         Examples:

 ₁H    1s^{1                           electronic configuration of hydrogen}

₂He   1s²                           electronic configuration of helium

₃Li     1s² 2s¹                   electronic configuration of lithum

₄Be   1s² 2s²                     electronic configuration of berylium

₅B     1s² 2s² 2p¹            electronic configuration of boron

₉F    1s² 2s² 2p⁵              electronic configuration of fluorine

₂₀Ca   1s² 2s² 2p⁶ 3s² 3p⁶ 4s²    electronic configuration of calcium

Periodic Table:

-         Is a table in which the elements arranged according to their atomic numbers.

-         Consist of 7 periods and 18 groups.

Position of elements in the modern periodic table        

     Determination of elements location in the periodic table:

The highest principle energy level represent  the element period and the

total number of electrons in this level represent the group of the element.

Ex:1

Then carbon is in the period two and group four

 

Ex:2

Then the Chlorine is in the period three and group seven

Periodic Table Properties (Properties and characteristics of atoms and elements):

In studying the properties of elements, we will stydy the trend of these properties through the periodic table

 

(1) - Atomic size or radius:

In the groups:

The atomic radius increase from top to bottom due to the increase in atomic number(increase the number of electrons in new main energy level)

In the periods:

The atomic radius decrease from left to right due to the increase of attraction of nucleus while the number of electrons increase in the same energy level.

 

Example:

The atomic radius of potassium is larger than the radius of sodium

₁₁Na    1s², 2s²  2p⁶  , 3s¹

₁₉K       1s², 2s²  2p⁶  , 3s²  3p⁶ 4s¹

The reason is that the increase of atomic number(number of electrons) and this increase is new principle energy level.

 

(2)– Ionization energy:

-         It is the energy required to remove electrons from atom in gaseous state.

-         Atoms can loss one, two, or three electrons so there are first, second and third ionization energies.

-         In periods:

The ionization energy increases from left to right because the atomic radius decrease and the attraction of nucleus to the outer electrons increase.

-         In groups:

Ionization energy decrease from top to bottom due to the increase in atomic radius and the decrease of attraction of nucleus to the outer electrons (this make elctrons need low energy to be removed).

-         Examples:

Sodium has only first ionization energy

                 ₁₁Na    1s², 2s² 2p6  , 3s¹                             Na⁺   1s², 2s²  2p⁶  

            Magnesium may have first or second ionization energies

           ₁₂Mg    1s², 2s² 2p6  , 3s²

                     Mg⁺    1s², 2s²  2p⁶  , 3s¹     (first ionization energy  of Mg)

                     Mg⁺²    1s², 2s²  2p⁶              (second ionization energy of Mg)

-         Explain why: the third ionization energy of Mg is much higher than the first and the second ionization energies(or why Mg⁺³ is not formed) ?

Answer:

Because the third ionization energy is the energy required to remove three electrons from the atom and in Mg the third elctron is removed from the complete 2p level which is near from the nucleus and has higher attraction from nucleus.

 

(3)-Electropositivity:

-         It is the ability of atom to form positive ion.

the more the electropositivity the more easily electrons removed.

-         In periods:

The electropositivity decrease in periods from left to right due to the decrease of atomic size and this made the loss of elecrtons difficult

-         In Groups:

The electropositivity increase in groups from top to bottom due to the increase of atomic size and this made the loss of elecrtons easy

-         Ex: Potasium K is more electropositive than  Na because K has atomic radius larger than Na so, K  loss electrons easier than Na

 

(4)-Electronaffinity:

-         It is the ability of atom to form negative ion

-         The more the electronaffinity the more easily ability of atom to form negative ion

-         In periods:

The electronaffinity increase from left to right due to the  decrease in the atomic radius

-         In groups:

The electronaffinity decrease from top to bottom due to the  increase in the atomic radius.

(5)-Electronegativity:

-         It is the ability of atom to attract the electons of the bond

-         In periods:

Increase from left to right due to the decrease in the atomic radius.

-         In groups:

Decrease from top to bottom due to the increase in the atomic radius.

 

(6)- Oxidation and Reduction:

-         Oxidation: is the process of loss of electrons.

-         Reduction: is the process of gain of electrons.

-         Oxidizing agent: is the substance that gain electrons (reduction).

-         Reducing agent: is the substance that loss electrons (oxidation).

-         In periods:

-         The reducing agent decrease from left to right due to the decrease in atomic radius

-         The Oxidizing  agent increase from left to right due to the decrease in atomic radius

-         In groups:

-         The reducing agent increase from top to bottom due to the increase in atomic radius

-         The Oxidizing agent decrease from top to bottom due to the increase in atomic radius

 

            Examples on Oxidation and reduction:

                    

How to determine the valency of group:

1-     The more electronegative atom(usually nonmetals) in the group gain electrons and has negative charge

2-     The less electronegative atom(usually metals) in the group loss electrons and has positive charge

3-     Then calculate the total charge of the group by the sum of the positive and negative charges

Examples:

 

(1)- NO₃

₇N  1s² 2s² 2p³

₈O  1s² 2s² 2p⁴

Oxygen is more electronegative than nitrogen so, oxygen gain electrons while nitrogen loss electrons;

We find that the three oxygen atoms need 6 electrons (-6) and nitrogen can loss only five electrons(+5) then +5 + (-6) = -1

N⁺⁵O^-6       then      NO₃^-

 

                               

           (2)- SO₄

   ₈O  1s² 2s² 2p⁴

     ₁₆S  1s² 2s² 2p⁴3s² 3p⁴

Oxygen is more electronegative than sulphur  so, oxygen gain electrons while sulphur loss electrons;

We find that the four oxygen atoms need 8 electrons(-8) but sulphur can loss only six electrons(+6) then +5 + (-8) = -2

                                               Then  SO₄^-2

 

(3)-CO₃

₈O  1s² 2s² 2p⁴

₆C  1s² 2s² 2p²

Oxygen is more electronegative than carbon so, oxygen gain electrons while carbon loss electrons;

We find that the three oxygen atoms need 6 electrons(-8) but carbon can loss only four electrons(+4) then +4 + (-6) = -2

                                              Then CO₃^-2

(8)-Valency and extended valency :

Valency: is the number of electrons lost or gained by the atom

Extended valency: is the valency result when atom take valency different from the normal valency

 

Ex:

H⁺¹O^-2Cl⁺¹      hybochlorous acid

H⁺¹Cl⁺⁵O₃^-6    Chloric acid

H⁺¹Cl⁺⁷O₄^-8    Perchloric acid

 

How to determine the oxidation state of element in a compound:

Useful rules:

-         Hydrogen=+1 when react with more electronegative atoms.

-         Hydrogen=-1 when react with more electronegative atoms (usually  group IA metals such as Na, K).

-         Group IA elements =+1, Group IIA elements =+2.

-         Group VIIA(Halogens)= -1 when react with more electronegative atoms

-         Oxygen usually =-2.

-         The oxidation state of any neutral compound=0

Examples:

1-     Determination of valency of Mn in KMnO₄

           K=+1 (group IA )

           O= -2  then 4O =4 x -2= -8

           (+1)+(-8)=+7

           Then K⁺¹Mn⁺⁷O₄^-8 then Mn =+7

 

2-     Determination of valency of Cr in K₂Cr₂O₇

K=+1 (group IA ) then 2 x +1= +2

O= -2  then 7O =7 x -2= -14

(+2) + (-14) =+12

Then K₂⁺²Cr₂⁺¹²O₇^-14   then  2Cr =+12

 Then  Cr=+6.

 

(9)-Acids and Bases:

There are three theories that explain the acids and bases as follow

 

(a)-Arrhenius theory:

This theory define the acid and base as follow:

Acid: is a substance that releases hydrogen ion when dissolved in water.

Ex: HCl, H₂So₄, H₃pO₄ CH₃COOH

Base: is a substance that releases hydroxyl ion when dissolved in water.

Ex: NaOH, KOH,and Ca(OH)₂

 

(b)-Bronsted theory

Acid: is a substance that loss proton (H⁺)

Base: is a substance that gain proton (H⁺)

(c)-Lewis theory:

Acid: is a substance that gains electrons (electron acceptor).

Ex: BCl₃, BeF₂ are acids because they have empty orbitals in which they can accept electrons as follow:

 

BCl_3:

₅B 1s² 2s² 2p¹  

₁₇Cl   1s²  2s²  2p⁶  3s²  3p⁵

  

We find that in Boron one electron is excited from 2s to 2p so we now have three orbitals with single electrons which are shared with 3Cl atoms .After that we find that the boron still has empty orbitals in which it can accept electrons, So BCl₃ act as Lewis acid.

 

BeF₂

₄Be  1s² 2s²

₉F  1s² 2s² 2p⁵

We find that Berillium one electron is excited from 2s to 2p so we now have two orbitals with single electrons which are shared with 2F atoms .After that we find that the Berillium still has empty orbitals in which it can accept electrons, So BeF₂ act as lewis acid.

 

base: is a substance that loss electrons(electron donor).

Ex: NH₃, PH₃ are Lewis bases because they have lone pair of electrons which they can donate to Lewis acid as follow:

 

NH₃

₇N  1s² 2s² 2p³

We find that nitrogen has five electrons in the outer shell, and it share with three hydrogen atoms by three electrons and after that the nitrogen will still has another two electrons which it can donate to any Lewis acid and forming coordinate covalent bond.

 

PH₃

₁₅P  1s² 2s² 2p⁶ 3s² 3p³

We find that phosphorus has five electrons in the outer shell, and it share with three hydrogen atoms by three electrons and after that the phosphorus will still has another two electrons which it can donate to any Lewis acid and forming coordinate covalent bond.

See Also:  Groups of Periodic Table and Trends in Periodic Table

See Also:  Types of Hydrogen (Isotopes of hydrogen)